Acid-Base Indicators in Titrations

Acid-Base Indicators

Definition and Role in Titrations

Acid-base indicators are weak acids or bases that exhibit different colors in their acid and base forms. They are used to signal the endpoint of a titration by changing color at a specific pH range.

• Endpoint: The point in a titration where the indicator changes color, ideally close to the equivalence point.

• Equivalence point: The point where the amount of titrant added is stoichiometrically equivalent to the amount of analyte in solution.

• Color change: Occurs when the ratio of the indicator's acid and base forms shifts as the pH passes through the indicator's transition range.

Indicator Color Change and pH

The color change interval for an indicator is typically centered around its pKa value. The general rule is:

• Indicator changes color when pH ≈ pKa ± 1.

For a weak acid indicator (HIn):

• At low pH: Color of HIn (acid form)

• At high pH: Color of In- (base form)

Table: Common Acid-Base Indicators

Indicator	pH Range	Colors
Crystal Violet	0.0–2.0	Violet to yellow
Methyl Orange	3.2–4.4	Red to yellow
Bromothymol Blue	6.0–7.6	Yellow to blue
Methyl Red	4.4–6.2	Red to yellow
Phenolphthalein	8.2–10.0	Colorless to pink

Indicator Selection for Titrations

To select the best indicator for a titration, choose one whose color change interval overlaps the pH at the equivalence point of the titration.

• Strong acid vs. strong base: Equivalence point at pH ≈ 7; use indicators like bromothymol blue.

• Strong acid vs. weak base: Equivalence point at pH < 7; use indicators with lower pKa.

• Weak acid vs. strong base: Equivalence point at pH > 7; use indicators with higher pKa.

Equations and Relationships

• The color change occurs when: $\text{pH} = \text{p}K_\text{a} \pm 1$

• For an indicator HIn: $\text{HIn} \rightleftharpoons \text{H}^+ + \text{In}^-$

• Henderson-Hasselbalch equation for indicators: $\text{pH} = \text{p}K_\text{a} + \log\left(\frac{[\text{In}^-]}{[\text{HIn}] }\right)$

Example: Choosing an Indicator

Given a titration curve with an equivalence point at pH ≈ 7, the best indicator would be one whose transition range includes pH 7, such as bromothymol blue (pKa ≈ 7.1).

Practice Problems

• Problem: A student uses an indicator with pKa = 4.7 for a strong acid–strong base titration. What is the pH range for the color change? Solution: pH = 4.7 ± 1 → pH range is 3.7 to 5.7.

• Problem: Bromophenol blue (pKa = 4.1) is yellow in acid and blue in base. If the solution has pH = 4.6, what color is observed? Solution: Since pH > pKa, the indicator is mostly in its basic (blue) form.

Summary Table: Indicator Properties

Indicator	pKa	Acid Color	Base Color	Transition Range
Bromophenol Blue	4.1	Yellow	Blue	3.1–5.1
Bromothymol Blue	7.1	Yellow	Blue	6.1–8.1
Phenolphthalein	9.4	Colorless	Pink	8.4–10.4

Additional info: The selection of an appropriate indicator is crucial for accurate titration results. The indicator should have a transition range that matches the expected pH at the equivalence point of the titration.